Are flashcards effective for AP Chemistry?

Yes. Flashcards use active recall — retrieving an answer from memory — which research shows is one of the most effective study methods, especially in short, spaced sessions. AP Chemistry rewards instant recall of definitions, formulas, reaction types, and trends, so flashcards make that knowledge automatic.

What topics do the AP Chemistry flashcards cover?

All nine units: atomic structure and periodic trends, bonding and VSEPR, intermolecular forces and solutions, reaction types and redox, kinetics and rate laws, thermodynamics and calorimetry, equilibrium and Le Chatelier, acids, bases, buffers and pH, and entropy, free energy, and electrochemistry.

How should I study with these AP Chemistry flashcards?

Lead with the heaviest units — Unit 3 (intermolecular forces) and Unit 8 (acids and bases), which together are about a third of the multiple-choice section. Mix the modes: flip to learn, type to test recall, match for speed, and quiz to check yourself before working full practice questions.

Are the AP Chemistry flashcards free?

Yes — 100% free, all four study modes, no paywall.

Do these flashcards match the current AP Chemistry course?

Yes. The cards are organized to the College Board's current AP Chemistry Course and Exam Description and its nine units, and reflect the two-section, 1–5-scored exam. Every card is built from the official content outline.

The SI unit for amount of substance; one mole = Avogadro's number of particles, $6.022\times10^{23}$.

$6.022\times10^{23}$ — the number of particles in one mole. Moles × this = number of particles.

Mass of one mole of a substance in g/mol, equal to the sum of the atomic masses in its formula.

Divide mass by molar mass: e.g. 88 g of $\ce{CO2}$ (44 g/mol) = 2 mol.

Multiply moles by Avogadro's number: 2 mol = $2\times6.022\times10^{23}=1.2\times10^{24}$ particles.

Separates isotopes by mass-to-charge ratio; peak positions give isotope masses and heights give relative abundances.

Atoms of the same element with different numbers of neutrons, hence different mass numbers.

The weighted average of an element's isotope masses by natural abundance; pulled toward the more abundant isotope.

Photoelectron spectroscopy (PES)

Measures energy to remove electrons; each peak = a subshell, peak height ∝ number of electrons in it.

PES peak height ratio

A filled s subshell (2 e⁻) vs a filled p subshell (6 e⁻) gives a 1-to-3 peak-height ratio.

Electron configuration

The arrangement of electrons by subshell in order of increasing energy, e.g. Mg = $\text{1s}^2\text{2s}^2\text{2p}^6\text{3s}^2$.

Electrons fill the lowest-energy orbitals first before higher ones.

Effective nuclear charge

The net positive pull on a valence electron after shielding by core electrons; drives periodic trends.

Inner core electrons block valence electrons from the full nuclear charge because they lie between them and the nucleus.

s electrons penetrate closer to the nucleus than p electrons in the same shell, so they are lower in energy.

Increases down a group and to the left across a period (larger as you move away from the top-right).

Energy to remove an electron from a gaseous atom; increases across a period, decreases down a group.

How strongly an atom attracts shared bonding electrons; increases up and to the right — fluorine is highest.

Coulomb's law (atomic)

Attraction increases with greater charge and decreases with greater distance — the basis of periodic trends.

A positively charged ion formed when an atom loses electrons; smaller than its parent atom.

A negatively charged ion formed when an atom gains electrons; larger than its parent atom.

The outermost-shell electrons that participate in bonding and determine chemical behavior.

Inner-shell electrons that shield valence electrons and do not participate in bonding.

The simplest whole-number ratio of atoms in a compound, found from percent composition by mass.

Electron transfer from a metal to a nonmetal forming oppositely charged ions; conducts when molten or dissolved.

A bond formed by sharing electron pairs between two nonmetals.

A 'sea' of delocalized electrons shared among metal cations; explains conductivity and malleability.

A diagram of a molecule's bonding and lone-pair electrons used to predict shape and reactivity.

Atoms tend to gain, lose, or share electrons to reach eight valence electrons (a noble-gas configuration).

An exception where the central atom has fewer than 8 electrons, e.g. boron in $\ce{BF3}$ (6 electrons).

An exception where a period-3+ central atom holds more than 8 electrons, e.g. $\ce{PCl5}$.

Electron domains around a central atom spread out to minimize repulsion, setting the molecular geometry.

A region of electron density — a bond (single, double, or triple) or a lone pair — counted in VSEPR.

2 electron domains, bond angle 180°, e.g. $\ce{CO2}$.

3 domains, 0 lone pairs, bond angle 120°, e.g. $\ce{BF3}$.

4 domains, 0 lone pairs, bond angle 109.5°, e.g. $\ce{CH4}$.

4 domains with 1 lone pair, angle ~107°, e.g. $\ce{NH3}$.

4 domains with 2 lone pairs, angle ~104.5°, e.g. $\ce{H2O}$.

5 bonding domains, angles 90° and 120°, e.g. $\ce{PCl5}$.

6 bonding domains, all 90° angles, e.g. $\ce{SF6}$.

5 domains with 1 lone pair (placed equatorial), e.g. $\ce{SF4}$.

6 domains with 2 lone pairs on opposite sides, e.g. $\ce{XeF4}$.

Head-on orbital overlap along the bond axis; every single bond is one sigma bond.

Side-by-side p-orbital overlap above and below the axis; double = 1 pi, triple = 2 pi.

The number of bonding electron pairs between two atoms; higher order = shorter, stronger bond.

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AP Chemistry Flashcards

Mole: The SI unit for amount of substance; one mole = Avogadro's number of particles, $6.022\times10^{23}$ .
Avogadro's number: $6.022\times10^{23}$ — the number of particles in one mole. Moles × this = number of particles.
Molar mass: Mass of one mole of a substance in g/mol, equal to the sum of the atomic masses in its formula.
Grams to moles: Divide mass by molar mass: e.g. 88 g of $\ce{CO2}$ (44 g/mol) = 2 mol.
Moles to particles: Multiply moles by Avogadro's number: 2 mol = $2\times6.022\times10^{23}=1.2\times10^{24}$ particles.
Mass spectrometry: Separates isotopes by mass-to-charge ratio; peak positions give isotope masses and heights give relative abundances.
Isotopes: Atoms of the same element with different numbers of neutrons, hence different mass numbers.
Average atomic mass: The weighted average of an element's isotope masses by natural abundance; pulled toward the more abundant isotope.
Photoelectron spectroscopy (PES): Measures energy to remove electrons; each peak = a subshell, peak height ∝ number of electrons in it.
PES peak height ratio: A filled s subshell (2 e⁻) vs a filled p subshell (6 e⁻) gives a 1-to-3 peak-height ratio.
Electron configuration: The arrangement of electrons by subshell in order of increasing energy, e.g. Mg = $\text{1s}^2\text{2s}^2\text{2p}^6\text{3s}^2$ .
Aufbau principle: Electrons fill the lowest-energy orbitals first before higher ones.
Effective nuclear charge: The net positive pull on a valence electron after shielding by core electrons; drives periodic trends.
Shielding: Inner core electrons block valence electrons from the full nuclear charge because they lie between them and the nucleus.
Penetration: s electrons penetrate closer to the nucleus than p electrons in the same shell, so they are lower in energy.
Atomic radius trend: Increases down a group and to the left across a period (larger as you move away from the top-right).
Ionization energy: Energy to remove an electron from a gaseous atom; increases across a period, decreases down a group.
Electronegativity: How strongly an atom attracts shared bonding electrons; increases up and to the right — fluorine is highest.
Coulomb's law (atomic): Attraction increases with greater charge and decreases with greater distance — the basis of periodic trends.
Cation: A positively charged ion formed when an atom loses electrons; smaller than its parent atom.
Anion: A negatively charged ion formed when an atom gains electrons; larger than its parent atom.
Valence electrons: The outermost-shell electrons that participate in bonding and determine chemical behavior.
Core electrons: Inner-shell electrons that shield valence electrons and do not participate in bonding.
Empirical formula: The simplest whole-number ratio of atoms in a compound, found from percent composition by mass.
Ionic bond: Electron transfer from a metal to a nonmetal forming oppositely charged ions; conducts when molten or dissolved.
Covalent bond: A bond formed by sharing electron pairs between two nonmetals.
Metallic bonding: A 'sea' of delocalized electrons shared among metal cations; explains conductivity and malleability.
Lewis structure: A diagram of a molecule's bonding and lone-pair electrons used to predict shape and reactivity.
Octet rule: Atoms tend to gain, lose, or share electrons to reach eight valence electrons (a noble-gas configuration).
Incomplete octet: An exception where the central atom has fewer than 8 electrons, e.g. boron in $\ce{BF3}$ (6 electrons).
Expanded octet: An exception where a period-3+ central atom holds more than 8 electrons, e.g. $\ce{PCl5}$ .
VSEPR theory: Electron domains around a central atom spread out to minimize repulsion, setting the molecular geometry.
Electron domain: A region of electron density — a bond (single, double, or triple) or a lone pair — counted in VSEPR.
Linear geometry: 2 electron domains, bond angle 180°, e.g. $\ce{CO2}$ .
Trigonal planar: 3 domains, 0 lone pairs, bond angle 120°, e.g. $\ce{BF3}$ .
Tetrahedral: 4 domains, 0 lone pairs, bond angle 109.5°, e.g. $\ce{CH4}$ .
Trigonal pyramidal: 4 domains with 1 lone pair, angle ~107°, e.g. $\ce{NH3}$ .
Bent geometry: 4 domains with 2 lone pairs, angle ~104.5°, e.g. $\ce{H2O}$ .
Trigonal bipyramidal: 5 bonding domains, angles 90° and 120°, e.g. $\ce{PCl5}$ .
Octahedral: 6 bonding domains, all 90° angles, e.g. $\ce{SF6}$ .
Seesaw geometry: 5 domains with 1 lone pair (placed equatorial), e.g. $\ce{SF4}$ .
Square planar: 6 domains with 2 lone pairs on opposite sides, e.g. $\ce{XeF4}$ .
Sigma (σ) bond: Head-on orbital overlap along the bond axis; every single bond is one sigma bond.
Pi (π) bond: Side-by-side p-orbital overlap above and below the axis; double = 1 pi, triple = 2 pi.
Bond order: The number of bonding electron pairs between two atoms; higher order = shorter, stronger bond.
Hybridization: Mixing of atomic orbitals: sp (linear), sp² (trigonal planar), sp³ (tetrahedral) to match geometry.
Bond polarity: Unequal electron sharing from an electronegativity difference, creating a bond dipole.
Molecular polarity: Overall dipole from bond dipoles and geometry; symmetric $\ce{CO2}$ is nonpolar, bent $\ce{H2O}$ is polar.
Intermolecular forces (IMFs): Attractions between molecules that set boiling point, vapor pressure, and solubility.
London dispersion forces: Weak attractions from temporary, instantaneous dipoles; present in all molecules.
Polarizability: How easily an electron cloud distorts; larger clouds (more electrons) give stronger dispersion forces.
Dipole-dipole forces: Attractions between the permanent dipoles of two polar molecules.
Hydrogen bonding: A strong IMF when H is bonded to F, O, or N; explains water's high boiling point.
Ion-dipole forces: Attraction between an ion and a polar molecule, e.g. $\ce{Na+}$ hydrated by water; strongest of the IMFs.
IMF strength order: Ion-dipole > hydrogen bonding > dipole-dipole > London dispersion.
IMF and boiling point: Stronger intermolecular forces require more energy to separate molecules, raising the boiling point.
Vapor pressure: Pressure of vapor in equilibrium with its liquid; lower when IMFs are stronger.
Volatile liquid: A liquid with weak IMFs, high vapor pressure, that evaporates readily.
Boiling: Occurs when a liquid's vapor pressure equals the external atmospheric pressure.
Branched vs straight chain: A branched isomer has less surface contact, weaker dispersion forces, and a lower boiling point.
Solid: A state with particles in fixed positions, definite shape and volume; strongest particle interactions.
Liquid: A state with definite volume but no fixed shape; particles touch but can move past one another.
Gas: A state with no fixed shape or volume; particles are far apart with negligible interactions.
Ideal gas law: $PV = nRT$ relates pressure, volume, moles, and temperature for an ideal gas.
Kinetic molecular theory: Gas particles are in constant random motion; average kinetic energy is proportional to absolute temperature.
Partial pressure: The pressure one gas in a mixture would exert alone; total pressure is the sum (Dalton's law).
Real vs ideal gas: Gases deviate from ideal behavior at high pressure and low temperature, where IMFs and molecular volume matter.
Solution: A homogeneous mixture of a solute dissolved in a solvent.
Like dissolves like: Polar solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes.
Molarity: Concentration in moles of solute per liter of solution: $M = \text{mol}/\text{L}$ .
Dilution: Adding solvent lowers concentration while moles of solute stay constant: $M_1V_1 = M_2V_2$ .
Miscible: Two liquids that mix in all proportions, like water and ethanol.
Distillation: Separating a mixture by differences in boiling point (vapor pressure).
Chromatography: Separating components by their differing attractions to a stationary vs a mobile phase.
Filtration: Separating an insoluble solid from a liquid by passing the mixture through a porous barrier.
Beer-Lambert law: Absorbance is proportional to concentration and path length: $A = \varepsilon b c$ .
Capillary action: Liquid rising in a narrow tube due to adhesion and cohesion (intermolecular forces).
Surface tension: Resistance of a liquid surface to increase its area, caused by cohesive intermolecular forces.
Viscosity: A liquid's resistance to flow; greater with stronger intermolecular forces.
Phase diagram: A plot of pressure vs temperature showing the solid, liquid, and gas regions and their boundaries.
Sublimation: A substance going directly from solid to gas without passing through the liquid phase.
Electrolyte: A solute that produces ions in solution and conducts electricity; strong electrolytes fully dissociate.
Chemical reaction: A process that rearranges atoms to form new substances, conserving mass and atoms.
Balanced equation: An equation with equal numbers of each atom on both sides; coefficients give the mole ratio.
Stoichiometry: Using the mole ratios in a balanced equation to relate amounts of reactants and products.
Limiting reactant: The reactant that runs out first and caps the amount of product formed.
Excess reactant: A reactant present in more than the stoichiometric amount; some is left over.
Percent yield: Actual yield ÷ theoretical yield × 100%, a measure of reaction efficiency.
Synthesis reaction: Two or more substances combine into one product (A + B → AB).
Decomposition reaction: One compound breaks down into two or more simpler substances.
Single-replacement reaction: One element replaces another in a compound; a redox reaction.
Double-replacement reaction: Ions exchange partners; often forms a precipitate, gas, or water.
Combustion reaction: A fuel reacts with oxygen, releasing heat; complete combustion gives $\ce{CO2}$ and $\ce{H2O}$ .
Incomplete combustion: Burning in limited oxygen produces carbon monoxide $\ce{CO}$ or soot instead of all $\ce{CO2}$ .
Precipitation reaction: Two solutions mix and form an insoluble solid (precipitate).
Spectator ion: An ion that appears unchanged on both sides of an equation and is canceled in the net ionic equation.
Complete ionic equation: An equation with all soluble strong electrolytes written as separate dissociated ions.
Net ionic equation: Shows only the species that actually react, after spectator ions are removed.
Oxidation: Loss of electrons; the oxidation number increases (OIL — Oxidation Is Loss).
Reduction: Gain of electrons; the oxidation number decreases (RIG — Reduction Is Gain).
Oxidation number: A bookkeeping charge assigned to an atom to track electron transfer in redox reactions.
Oxidizing agent: The species that is reduced (gains electrons) and causes another to be oxidized.
Reducing agent: The species that is oxidized (loses electrons) and causes another to be reduced.
Titration: Adding a known concentration of titrant until a reaction is complete to find an unknown amount.
Reaction rate: How fast reactants are consumed or products formed; change in concentration over time.
Average rate: Change in concentration of a species divided by the elapsed time interval.
Rate law: $\text{rate}=k[A]^x[B]^y$ ; the orders x and y come from experiment, not coefficients.
Rate constant (k): The proportionality factor in a rate law; depends on temperature and activation energy, not concentration.
Reaction order: The exponent on a reactant in the rate law; the overall order is the sum of the exponents.
Zero order: Rate is independent of concentration; a plot of [A] vs time is a straight line.
First order: Rate ∝ [A]; a plot of ln[A] vs time is linear; units of k are $\text{s}^{-1}$ .
Second order: Rate ∝ [A]²; a plot of $1/[A]$ vs time is linear.
Method of initial rates: Find order by changing one reactant: no change = 0, doubles = 1, quadruples (×4) = 2.
Half-life: Time for a reactant's concentration to fall by half; constant for first-order reactions.
Collision theory: Reactions occur when particles collide with enough energy and correct orientation.
Activation energy (Eₐ): The minimum energy needed to react — the barrier between reactants and the transition state.
Transition state: The high-energy 'activated complex' at the peak of the energy profile.
Arrhenius equation: Relates the rate constant to temperature and activation energy; higher T or lower Eₐ speeds the reaction.
Catalyst: A substance that speeds a reaction via a lower-Eₐ pathway and is not consumed.
Reaction mechanism: The step-by-step sequence of elementary steps by which a reaction occurs.
Rate-determining step: The slowest elementary step; it limits the overall reaction rate.
Intermediate: A species produced in one step and consumed in a later step; not in the overall equation.
Temperature and rate: Raising temperature increases the fraction of collisions with energy above Eₐ, speeding the reaction.
Catalyst and equilibrium: A catalyst speeds both directions equally; it does not shift equilibrium or change K.
System: The specific part of the universe being studied, such as the reacting chemicals.
Surroundings: Everything outside the system that can exchange energy with it.
Heat (q): Energy transferred because of a temperature difference; flows from warmer to cooler.
Temperature: A measure of the average kinetic energy of particles, distinct from total heat.
Endothermic: A process that absorbs heat; $q$ and $\Delta H$ are positive (feels cold).
Exothermic: A process that releases heat; $q$ and $\Delta H$ are negative (feels hot).
Enthalpy (H): Heat content at constant pressure; $\Delta H$ is the heat absorbed or released.
ΔH sign convention: $\Delta H < 0$ = exothermic (products lower in energy); $\Delta H > 0$ = endothermic.
Specific heat capacity: Energy to raise 1 g of a substance by 1 °C; water's is unusually high.
Heat capacity: Energy to raise an entire object/sample by 1 °C (no per-gram basis).
Calorimetry equation: $q = mc\Delta T$ : heat = mass × specific heat × temperature change.
Calorimeter: An insulated device that measures heat flow by tracking temperature change.
Enthalpy of solution: The heat absorbed or released when a solute dissolves in a solvent.
Hess's law: The $\Delta H$ of an overall reaction equals the sum of the step $\Delta H$ values (state function).
State function: A property (like enthalpy) that depends only on the current state, not the path taken.
Bond enthalpy: Energy to break a bond; $\Delta H$ ≈ bonds broken minus bonds formed.
Standard enthalpy of formation: $\Delta H_f^\circ$ : heat change forming 1 mol of a compound from its elements in standard states.
Thermal equilibrium: The state where two objects in contact reach the same temperature and net heat flow stops.
Chemical equilibrium: A dynamic state where forward and reverse reaction rates are equal and concentrations stay constant.
Equilibrium constant (K): Ratio of product to reactant concentrations at equilibrium, each raised to its coefficient.
Large vs small K: A large K (≫1) favors products; a small K (≪1) favors reactants.
Reaction quotient (Q): The same ratio as K but at any moment; comparing Q to K predicts the shift.
Q vs K: $Q<K$ shifts forward; $Q>K$ shifts reverse; $Q=K$ is at equilibrium.
Kc vs Kp: Kc uses concentrations, Kp uses partial pressures; $K_p = K_c(RT)^{\Delta n}$ .
Equilibria and pure phases: Pure solids and liquids are left out of the K expression (their activity is 1).
Le Chatelier's principle: A system at equilibrium shifts to partially offset an applied stress.
Adding reactant: Shifts equilibrium toward products to consume the added reactant.
Removing product: Shifts equilibrium toward products to replace what was removed.
Pressure/volume change: Increasing pressure (lower volume) shifts toward the side with fewer moles of gas.
Temperature and K: Only a temperature change actually changes the value of K; heating an exothermic reaction lowers K.
Catalyst at equilibrium: A catalyst reaches equilibrium faster but does not shift it or change K.
Solubility product (Ksp): Equilibrium constant for a slightly soluble salt dissolving into its ions.
Molar solubility: Moles of a salt that dissolve per liter; for a 1:1 salt, $K_{sp}=s^2$ .
Predicting precipitation: A precipitate forms when the ion product Q exceeds Ksp.
Common-ion effect: Adding an ion already in a solubility equilibrium decreases the salt's solubility.
ICE table: A Reactants/Initial-Change-Equilibrium chart used to solve equilibrium problems.
Arrhenius acid: A substance that produces $\ce{H+}$ (hydronium) ions in water.
Arrhenius base: A substance that produces $\ce{OH-}$ (hydroxide) ions in water.
Bronsted-Lowry acid: A proton ( $\ce{H+}$ ) donor.
Bronsted-Lowry base: A proton ( $\ce{H+}$ ) acceptor.
Conjugate acid-base pair: Two species differing by one proton, e.g. $\ce{NH4+}$ and $\ce{NH3}$ .
Strong vs weak acid: A strong acid fully ionizes; a weak acid only partially ionizes, setting up an equilibrium.
Strong acid–weak conjugate: The stronger the acid, the weaker its conjugate base, which has little tendency to reaccept the proton.
Amphoteric: A substance that can act as either an acid or a base, like water.
pH: $\text{pH} = -\log[\ce{H+}]$ ; below 7 acidic, 7 neutral, above 7 basic.
pOH: $\text{pOH} = -\log[\ce{OH-}]$ ; at 25 °C, pH + pOH = 14.
Kw: The ion-product of water, $K_w = [\ce{H+}][\ce{OH-}] = 1.0\times10^{-14}$ at 25 °C.
pH of a strong acid: Equals $-\log$ of its concentration, since a strong acid fully ionizes.
Ka: Acid dissociation constant; a larger Ka (smaller pKa) means a stronger acid.
pKa: $\text{p}K_a = -\log K_a$ ; a lower pKa means a stronger acid.
Buffer: A weak acid and its conjugate base together; resists pH change on adding small amounts of acid or base.
Henderson-Hasselbalch: $\text{pH} = \text{p}K_a + \log\frac{[\text{base}]}{[\text{acid}]}$ .
Buffer at pH = pKa: When [base] = [acid] (1:1 ratio), the log term is 0, so pH = pKa.
Buffer capacity: The amount of acid or base a buffer can neutralize before pH changes significantly.
Equivalence point: The point in a titration where moles of titrant exactly neutralize the analyte.
Buffer region: The region before the equivalence point where weak acid and conjugate base coexist and pH changes slowly.
Weak acid–strong base titration: Equivalence point pH is above 7 because the conjugate base is basic.
Indicator: A substance that changes color over a pH range; choose one whose range spans the equivalence point.
Phenolphthalein: An indicator colorless below ~8 and pink above ~10; good for weak-acid/strong-base titrations.
Neutralization: An acid and base react to form water and a salt; $\ce{H+ + OH- -> H2O}$ .
Entropy (S): A measure of the dispersal of matter and energy; the number of accessible microstates.
Entropy and gases: Forming more moles of gas increases disorder, so $\Delta S$ is positive.
Third law of thermodynamics: A perfect crystalline solid at absolute zero (0 K) has zero entropy.
Gibbs free energy: $\Delta G = \Delta H - T\Delta S$ ; the criterion for thermodynamic favorability.
Thermodynamically favorable: A process with $\Delta G < 0$ ; favorable does not mean fast.
Enthalpy-driven: A favorable process driven mainly by a large negative $\Delta H$ .
Entropy-driven: A favorable process driven mainly by a large positive $\Delta S$ .
Crossover temperature: Where favorability flips; set $\Delta G = 0$ , so $T = \Delta H/\Delta S$ .
ΔG and K: $\Delta G^\circ = -RT\ln K$ ; a negative $\Delta G^\circ$ corresponds to $K>1$ .
Galvanic (voltaic) cell: Uses a spontaneous redox reaction to produce electricity.
Anode: The electrode where oxidation occurs (loss of electrons).
Cathode: The electrode where reduction occurs (gain of electrons).
Standard hydrogen electrode: The reference half-cell defined as exactly $0$ V.
Standard cell potential: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ ; positive means spontaneous.
ΔG and cell potential: $\Delta G^\circ = -nFE^\circ$ ; a larger positive $E^\circ$ gives a more negative $\Delta G^\circ$ .
Electrolytic cell: Uses an external power source to drive a non-spontaneous redox reaction.

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By Career EmployerUpdated Sep 12, 2026

200+ free AP Chemistry flashcards — 4 ways to study

AP Chemistry Flashcard of the DayNew daily

The classic card. Do you know this one?

Click Study Flashcards above to open the flashcard hub — 200 AP Chemistry cards you can flip, match, type, or quiz yourself on. Every card is drawn from the College Board’s nine AP Chemistry units, so you study exactly what the exam measures.^[1] Pair them with our free practice test and study guide.

AP Chemistry is one of the 17 AP exams — explore our AP flashcards to compare and prep across the whole family.

AP Chemistry Flashcard Study Modes

Flip mode walks you through cards one at a time so you can check yourself. Match is a timed game that pairs terms with their definitions. Type shows the definition and asks you to produce the term, so a card like Molarity has to come back from memory and be spelled correctly. Quiz turns the same cards into multiple-choice questions.

Free AP Chemistry flashcards from Career Employer — active recall for all nine College Board units

Why Flashcards Work for AP Chemistry

Unit 3: Properties of Substances & Mixtures is the largest section at 34 cards, drilling phase and solution language side by side: Gas, Solid and Liquid sit next to Molarity, Dilution and Miscible, so you get both particle-level descriptions and the vocabulary that shows up in lab-style questions. Unit 1: Atomic Structure & Properties adds 24 cards on counting and electron arrangement, including Mole, Shielding and Core electrons.

Unit 2: Compound Structure & Properties holds 24 cards on bonding and molecular shape, with Ionic bond, VSEPR theory and Tetrahedral among the fronts you have to define precisely. Unit 8: Acids & Bases matches that at 24 cards and leans on the symbol set you need to read instantly, including Ka, pOH and Amphoteric.

Unit 4: Chemical Reactions covers 22 cards on reaction bookkeeping and electron transfer, including Stoichiometry, Percent yield and Oxidizing agent. Unit 5: Kinetics follows with 20 cards of rate vocabulary, where Rate law, Zero order and Intermediate keep the reaction orders separate from the mechanism terms.

Unit 6: Thermodynamics carries 18 cards on energy accounting, such as Heat (q), Hess’s law and Surroundings, where the sign conventions live. Unit 7: Equilibrium has 18 more built around comparisons and setups, including Q vs K, ICE table and Molar solubility, alongside shift prompts such as Removing product.

Unit 9: Applications of Thermodynamics closes the deck with 16 cards tying spontaneity to electrochemistry, so Gibbs free energy and ΔG and K appear next to cell vocabulary like Anode and Cathode.

AP Chemistry rewards instant recall of definitions, formulas, reaction types, and periodic and IMF trends.^[2] Spaced flashcards are the most efficient way to make that knowledge automatic. Used alongside our practice test and study guide, they turn review time into measurable progress.

AP Chemistry Flashcards by Unit

The cards are organized by the College Board’s nine AP Chemistry units. Drill the heaviest ones first — Unit 3 (intermolecular forces) and Unit 8 (acids and bases) together are about a third of the multiple-choice section:^[1]

AP Chemistry flashcards by College Board unit
Unit	Content
Unit 1	Atomic Structure & Properties — moles, PES, periodic trends
Unit 2	Compound Structure & Properties — bonding, VSEPR, hybridization
Unit 3	Properties of Substances & Mixtures — intermolecular forces, solutions
Unit 4	Chemical Reactions — reaction types, net ionic equations, redox
Unit 5	Kinetics — rate laws, mechanisms, activation energy
Unit 6	Thermodynamics — heat, enthalpy, calorimetry, Hess's law
Unit 7	Equilibrium — K and Q, Le Chatelier, solubility
Unit 8	Acids & Bases — pH, buffers, titrations
Unit 9	Applications of Thermodynamics — entropy, free energy, electrochemistry

How to Get the Most Out of These Flashcards

Start with the biggest block. Unit 3: Properties of Substances & Mixtures carries 34 cards, more than any other unit, so clearing it early removes the largest chunk of unfamiliar terms.
Type-drill the definitions you confuse. Cards like Molarity and Percent yield reward exact wording, and typing forces you to state the ratio instead of nodding along in Flip mode.
Use Match for short symbol fronts. The Unit 8: Acids & Bases cards such as Ka and pOH are ideal, since speed pairing trains the instant recognition those questions demand.
Move to the practice test once recall is automatic. When Unit 7: Equilibrium fronts like ICE table and Q vs K come back without hesitation, shift to full questions and the study guide.
Keep a steady cadence. Two units per sitting across 200 cards, then a mixed Quiz pass over units already seen, keeps older material like Unit 5: Kinetics from fading.

AP Chemistry Flashcards FAQ

Two hundred free AP Chemistry flashcards, organized across all nine College Board units — from Atomic Structure and Compound Structure through Kinetics, Equilibrium, Acids and Bases, and Applications of Thermodynamics. They're free with no account required.

References

1.College Board. “AP Chemistry Course and Exam Description.” College Board. ↑
2.College Board. “AP Chemistry Exam — AP Students.” College Board. ↑
3.College Board. “AP Chemistry Course Overview.” College Board. ↑

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